Sulphurous Acid

Formation

Sulphur dioxide imparts to water acidic properties, owing to the formation of an unstable sulphurous acid; the chemical properties of moist or dissolved sulphur dioxide are therefore very largely the properties of sulphurous acid.

By analogy with the corresponding salts the acid should possess a formula H2SO3, but the substance has not been isolated, the solutions decomposing and yielding free sulphur dioxide even when allowed to evaporate at the ordinary temperature. A number of solid products described as hydrates of the acid have been obtained; they are best produced by the prolonged passage of sulphur dioxide into water at about 0° C., when a crystalline deposit slowly forms. The products are easily dissociable, denser than water, and have melting-points not far removed from 4° C. It has been shown, however, that only one hydrate, H2SO3.5H2O, exists, and that the other products are to be regarded as mixtures of this hydrate with ice.

It is worthy of note that liquid sulphur dioxide and water are not completely miscible at low temperatures.

Properties of Sulphurous Acid

Sulphur dioxide dissolves in water with the evolution of heat. The aqueous solutions are denser than water. The ready decomposition of sulphurous acid is obvious from the pungent smell of sulphur dioxide always observable above the solution.

The acid is rather feeble, but it reddens litmus, decomposes carbonates and neutralises alkalis with the formation of salts. The baser metals are also slowly attacked, and the liquid finds frequent use as an etching reagent in the metallographie examination of alloys.

Although the solution conducts the electric current, the dissolved sulphur dioxide is mainly in an un-ionised condition and even unhydrated; what electrolytic dissociation does occur is mainly into the ions H• and HSO3', the proportion of SO3' ions being very small indeed.

The absorption spectra of the aqueous solutions indicate that the sulphur dioxide is mainly present as SO2 molecules, some hydrate molecules also being present, but only small amounts of H2SO3 and its ions.

Autoxidation

Aqueous solutions of sulphurous acid are well known to undergo slow autoxidation in the absence of air or oxidising agents, the change being usually represented by the equation:

3H2SO3 = 2H2SO4 + S + H2O.

accelerated by rise of temperature and by exposure to been suggested that the following initial reactions are

It is pointed out that although aqueous sulphurous acid appears to contain most of its sulphur dioxide as such, yet solutions of the alkali bisulphites appear to contain some pyrosulphite (metabisulphite). It is therefore not unreasonable to assume that some free pyrosulphurous acid is present in aqueous sulphurous acid solutions. Since hydrosulphurous acid is unstable under the prevailing conditions, it does not accumulate, but rapidly undergoes further change, which makes proof of its formation difficult. It is assumed that some of the colourless acid formed in reaction (1) undergoes rearrangement to the isomeric coloured form:

HO.S.O.SO.OH (colourless) ⇔ (HO)2S.SO2 (coloured) (2),

and that these two isomerides react, yielding pyrosulphurous and thiosulphuric acids, thus:

2H2S2O4 ⇔ H2S2O5 + H2S2O3 (3).

The thiosulphuric acid now undergoes further change, one of the first products being trithionic acid, H2S3O6, tetra- and penta-thionic acids being gradually formed as the concentration of sulphurous acid diminishes with the progress of autoxidation. These acids in their turn gradually decompose, the final products being sulphuric acid and free sulphur. This explains why thionic acids are almost always present in old solutions of sulphurous acid, whether these have been entirely protected from atmospheric oxidation or not, and the amount of these acids present will depend largely upon the age and composition of the solution.

It must be stated, however, that the absorption spectra of aqueous alkali hydrogen sulphite solutions are similar to those of sulphur dioxide, and even aqueous or alcohol solutions of potassium or ammonium pyrosulphite show no spectrum characteristic of the ion S2O5'. On warming or on exposure to light, such solutions show an increased absorption, the bands being displaced towards the longer wave-lengths, and this is to be explained by the photo-oxidation of HSO3' ions to SO4', with formation of SO2-hydrate.

Dilute solutions of sulphurous acid decompose more quickly and completely than concentrated solutions, while in presence of a strong acid the decomposition is inhibited, being completely suppressed in twice-normal hydrochloric acid solution.

It has been found that sulphurous acid liberated in solution from its salts at temperatures of 100° to 120° C. or lower, may undergo instantaneous auto-reduction with production of hydrogen sulphide:

4H2SO3 = H2S + 3H2SO4,

this corresponding with the change which takes place when solid sodium sulphite is heated above 150° C.:

4Na2SO2 = Na2S + 3Na2SO4.

It appears, moreover, that this rapid auto-reduction occurs only with sulphurous acid liberated from its salts and not with gaseous sulphur dioxide or saturated solutions of the gas or even with the solid hydrate. It seems possible that the high reducing power of the sulphurous acid under these conditions may be due to the momentary existence of the acid H.SO2.OH, produced from a corresponding metallic salt such as was postulated by Divers and Shimidzu.

A mixture of sodium sulphite and sodium hydrogen sulphite in aqueous solution at 150° C. undergoes a change in accordance with the equation:

2NaHSO3 + Na2SO3 = 2Na2SO4 + S + H2O.

Electrolysis of sulphurous acid causes anodic oxidation to sulphuric acid, whilst hydrosulphurous acid, H2S2O4, is produced by the reducing action of the current at the cathode; the former change may be rendered quantitative by the addition of a manganese salt; the hydrosulphurous acid may be accompanied by various of its decomposition products.

The most marked group of reactions exhibited by sulphurous acid is dependent on its reducing power. In dilute aqueous solution slow oxidation by free oxygen occurs with liberation of 63.6 Cals. per gram- molecule oxidised; the change is accelerated by the influence of light and the presence of salts of manganese, cobalt or iron. Sulphurous acid solution reduces selenious and tellurous acids to selenium and tellurium, respectively; it precipitates the noble metals, for example gold, from solutions of their salts. Salts derived from higher oxides are generally reduced to derivatives of lower oxides; thus ferric salts, chromates, manganates and permanganates yield ferrous, chromic and manganous salts, respectively.

In the case of the reduction of ferric salts by sulphur dioxide it has been suggested that the reduction most probably proceeds in stages, a red ferric ferrisulphite, Fe[Fe(SO3)3], being first formed, which on warming yields the ferrous salt and the dithionate, thus:

Fe[Fe(SO3)3] = FeS2O6 + FeSO3.

Oh boiling, the dithionate decomposes:

FeS2O6 + H2O = FeSO4 + H2SO3.

The halogen elements are converted into the corresponding hydracids by sulphurous acid, the change being reversible in the case of bromine and especially with iodine:

H2SO3 + I2 + H2O ⇔ H2SO4 + 2HI.

In this reaction it has been shown that there is an intermediate formation of a yellow compound of composition SO2.HI.

In an analogous manner cyanogen, which in many of its reactions exhibits a marked similarity to chlorine, is slowly reduced to hydrocyanic acid:

With the exception of perchloric acid, the halogen oxyacids are reduced through the corresponding halogen elements to the hydracids; in the case of iodic acid the formation of free iodine after a definite interval forms a striking example of a " time reaction." The formation of the element is not observed until after the disappearance of the whole of the sulphite. The mechanism of the reaction is probably as follows:

IO3' + 5I' + 6H• = 3I2 + 3H2O, SO3' + I2 + 2OH' = SO4' + 2I' + H2O,

the latter reaction being much more rapid than the former.

Nitric acid undergoes reduction by sulphurous acid less readily than nitrous acid, the product in each case being sulphuric acid with nitrous or nitric oxide: it is possible that nitrogen-sulphur acids such as nitrosulphonic acid are intermediately produced. Nitric oxide can be reduced slowly to nitrous oxide by sulphurous acid, but no further.

Mercuric chloride readily oxidises sulphurous acid, the reaction being quantitative if the solution of the latter is dilute:

H2SO8 + 2HgCl2 + H2O = H2SO4 + Hg2Cl2 + 2HCl.

The result is different if sodium hydrogen sulphite is used instead of sulphurous acid; the addition of excess of mercuric chloride then gives rise to an equimolecular proportion of hydrochloric acid, the titration of which provides a convenient method for estimating the quantity of hydrogen sulphite originally present:

2NaHSO3 + 2HgCl2 = 2HgCl.SO3Na + 2HCl.

By taking a sulphurous acid solution and titrating with sodium hydroxide and methyl orange until the sodium hydrogen sulphite is formed, then adding excess of mercuric chloride and compleyng the titration, sulphurous acid itself may be estimated.

The compound obtained by the interaction of mercuric chloride and ammonia dissolves when sulphur dioxide is passed through the liquid.

Sulphurous acid can, on the other hand, undergo reduction. Hydrogen sulphide reduces it, producing sulphur chiefly. The primary action may be represented by the reversible equation:

H2SO3 + H2SH2SO2 + H2SO.

The compound H2SO, which is assumed to be of the peroxide type, decomposes with precipitation of sulphur. The hydrogen sulphide also reacts further with the sulphoxylic acid primarily formed to yield more H2SO:

H2SO2 + H2S = 2H2SO,

so that the final result of the interaction is

2H2S + SO2 = 3S + 2H2O.

This final stage is reached more quickly in the presence of excess of hydrogen sulphide. In the presence of excess of sulphurous acid some polythionic acid is formed.

Phosphorus at 200° C. in a sealed tube gives phosphoric acid and hydrogen sulphide. Nascent hydrogen, for example from zinc and sulphuric acid, yields hydrogen sulphide, as also do phosphorous acid, titanous chloride, and even stannous chloride:

Formic acid gives a reducing action of an unusual type with sulphurous acid; the mixture of the two acids is a much stronger reducing agent than either of the components by itself, on account of the formation of hydrosulphurous acid:

H.CO2H + 2H2SO3 = H2S2O4 + 2H2O + CO2.

Hydrosulphurous acid is also formed by the action of many metals on sulphurous acid; the crust of ferrous sulphite and hydrosulphite formed on the surface of metallic iron is of great protective value in the iron cylinders used for storing liquid sulphur dioxide, which generally contains traces of water.

Detection and Estimation

Sulphurous acid is usually detected by its reducing action, for instance on potassium dichromate solution, the acid being warmed in order that the test may be effected with the evolved gas; the odour of sulphur dioxide is also a fairly trustworthy indication. For special purposes many of the reactions already mentioned may be applied.

A sensitive colour test for sulphite ions consists in adding, drop by drop, a 0.01 per cent, solution of Fast Blue R crystals, shaking after each addition, until the violet coloration disappears and a yellow solution is produced; the test is sensitive to one part of sulphurous acid in about 175,000. Thiosulphates and polythionates do not interfere, but sulphides and hydroxides must be absent.

If to a neutral solution of an alkali sulphite containing phenolphthalein a few drops of a 1 per cent, formaldehyde solution are added, a pink colour due to liberation of alkali hydroxide is produced:

H.CHO + Na2SO3 + H2O = H.CH(OH).SO3Na + NaOH.

When a solution containing sulphite is added to a few drops of Bettendorf's reagent (stannous chloride in concentrated hydrochloric acid) a yellowish-brown deposit of tin sulphides is formed, owing to reduction of the sulphite; 0.6 per cent, of the latter can thus be detected. Contrary to what might be expected, the estimation of sulphurous acid cannot be satisfactorily effected by direct titration with potassium permanganate in the presence of an acid, because, under ordinary conditions, a considerable proportion of the sulphurous acid escapes complete oxidation, being converted into dithionic acid. More satisfactory results are obtained by the addition of a large excess of permanganate (or dichromate) with sulphuric acid and the subsequent determination of the excess by a suitable standard solution, e.g. of oxalic acid or ferrous sulphate. Good results have been obtained by observing the following definite conditions: A known volume of acidified standard potassium permanganate solution is divided into two portions, one being reserved for comparison. To the other portion is added the sulphurous acid to be estimated, and more standard permanganate is added until the colour is again of the same intensity as that of the reserved portion. The solutions are mixed, divided into two portions and the process again repeated. These repetition operations ' are necessary owing to the fact that the colour obtained on re-titrating a mixture of sulphurous acid and potassium permanganate is slightly different in tint from that of the permanganate alone. The quantity of permanganate solution required to restore the standard colour is equivalent to the amount of sulphurous acid present.

Another method which is applicable also to the estimation of solutions containing sulphites and hydrogen sulphites is the addition of a known excess of acidified hydrogen peroxide to the solution of sulphurous acid. The excess of peroxide may then be titrated with standard permanganate solution. The reactions are as follows:

The more usual volumetric method for the estimation of sulphurous acid is the excess iodine method already described. The reaction between the acid and iodine is not reversible in the state of dilution obtaining in volumetric analysis, therefore the addition of sodium hydrogen carbonate, as often recommended, is unnecessary. The low results obtained when the sulphurous acid is exposed to air during the titration are due entirely to evaporation of sulphur dioxide, the amount of atmospheric oxidation being negligible. Sulphite solutions, on the other hand, readily undergo atmospheric oxidation, so that in order to obtain an accurate estimation of sulphite it is advisable to dissolve the salt, in a 5 per cent, recently boiled glycerol solution in a flask filled with carbon dioxide and then run in a measured excess of standard iodine solution also under carbon dioxide and titrate the residual iodine with thiosulphate. Satisfactory results may also be obtained by employing more powerful oxidising agents, such as an acid solution of potassium iodate, or sodium hypochlorite. In the latter case, excess of a standard hypochlorite solution is added to the sulphurous acid solution and titrated back with iodine. This method is applicable to solutions of high SO2-concentration.

It is also possible with sulphurous acid solutions, using methyl orange as indicator, to titrate with a standard solution of an alkali to the halfway "bisulphite" stage. If necessary the alkali hydrogen sulphite may then be estimated by the addition of mercuric chloride and further titration with alkali, as already described. By this double titration method it is possible to estimate sulphurous acid in the presence of other sulphur acids.

Occasionally it is desired to determine sulphurous acid by a gravimetric method; it is then usual to expel the sulphur dioxide from the solution under examination by distilling in an atmosphere of carbon dioxide and oxidise the gas to sulphuric acid by absorption in bromine water or iodine solution, subsequently adding barium chloride and weighing the precipitated barium sulphate. This method also gives accurate results volumetrically if steps are taken to prevent loss of iodine by volatilisation in the current of carbon dioxide. The excess iodine is titrated with a solution of sodium thiosulphate.

Constitution of Sulphurous Acid

In deciding the constitution of sulphurous acid, choice has to be made between the two alternatives and , which may respectively be designated as the unsymmetrical structure and the symmetrical structure. It may be stated at once that no absolute decision is at present possible with respect to the inorganic salts of sulphurous acid and the aqueous acid itself, although clear evidence is obtained with the organic sulphites.

When an alkali sulphite is allowed to react with the bromide or iodide of an organic radical, for example with ethyl iodide, the resulting compound, ethylsulphonic acid, undoubtedly possesses the unsymmetrical structure, , because not only does the corresponding diethyl compound (ethyl ethylsulphonate) merely undergo decomposition to the stage of the potassium salt C2H5.SO2.OK when heated with aqueous potassium hydroxide solution, but also the parent ethylsulphonic acid can be obtained by the oxidation of ethyl hydrogen sulphide, C2H5.S.H, in which the ethyl radical is certainly attached directly to sulphur.

On the other hand, thionyl chloride reacts with ethyl alcohol to produce a compound of the symmetrical structure . This diethyl sulphite is a colourless liquid of b.pt. 161° C., and is quite distinct from the foregoing isomeric ethyl ethylsulphonate (which boils at 207° C.), being decomposed by aqueous alkali with formation of alkali sulphite and ethyl alcohol and reacting with organo-magnesium compounds with the formation of a sulphoxide, , where R represents the organic radical.

Organic compounds are therefore obtainable representative of both the symmetrical and the unsymmetrical structures. Only one series of inorganic salts is known, however, and the task of definitely assigning to them one of the structures is a matter of great difficulty, because the purely inorganic evidence is conflicting.

Whilst the formation of thionyl chloride by the action of phosphorus pentachloride on an alkali sulphite favours a symmetrical structure for the latter, the interaction of an alkali sulphite with ethyl iodide, as mentioned already, is directly opposed to this evidence, as also is the production of sodium sulphite on reducing sodium dithionate with sodium. The remaining inorganic evidence is little more satisfactory. What appeared to be final evidence was once brought forward in the description of two isomeric salts obtained by the neutralisation of sodium hydrogen sulphite with potassium carbonate solution and of potassium hydrogen sulphite with sodium carbonate solution. These two salts dissolved in water to give identical solutions, in accordance with the theory of electrolytic dissociation, and their structures were assumed to be and , respectively, the former yielding crystals with two molecules of water of crystallisation, the latter with one. Subsequent workers, however, have been unable to confirm these observations.

No final or absolute decision is therefore yet possible. By analogy with other cases of a similar type amongst organic compounds, the possibility of a "dynamic isomerism" must not be left out of consideration. It is quite possible that in sulphurous acid solution, molecules of the symmetrical and the unsymmetrical constitution may be present, side by side, in equilibrium with one another, and that even in solutions of the salts a similar condition of equilibrium may exist. Indeed, the distinctly contradictory nature of some of the chemical evidence favours this view.